1. Chapter 10
Liquids and solids

2. 10.1 Intermolecular forces
Inside molecules (intramolecular) the atoms are bonded to each other.
Intermolecular refers to the forces between the molecules.
These are what hold the molecules together in the condensed states.

3. Intermolecular forces
Strong
covalent bonding
ionic bonding
Weak
Dipole dipole
London dispersion forces
Hydrogen bonding
During phase changes the molecules stay intact.
Energy used to overcome forces.

4. Dipole - Dipole Remember where the polar definition came from?
Molecules line up in the presence of a electric field. The opposite ends of the dipole can attract each other so the molecules stay close together.
1% as strong as covalent or ionic bonds.
Weaker with greater distance.
Small role in gases.

5. + - + -

6. Hydrogen Bonding
Especially strong dipole-dipole forces when H is attached to F, O, or N
These three because-
They have high electronegativity.
They are small enough to get close.
Affects boiling point.

7. Water
Hydrogen Bonding Clip d+ d- d+

8. Boiling Points 100 H2O HF 0ºC H2Te H2Se NH3 SbH3 H2S HI AsH3 HCl HBr
PH3
NH3
SbH3
AsH3
CH4
SiH4
GeH4
SnH4
-100
200

9. London Dispersion Forces
Non - polar molecules also exert forces on each other.
Otherwise, no solids or liquids.
Electrons are not evenly distributed at every instant in time.
Have an instantaneous dipole.
Induces a dipole in the atom next to it.
Induced dipole - induced dipole interaction.

10. Example H d+ d d- H d+ d- H
LD Video

11. London Dispersion Forces
Weak, short lived.
Lasts longer at low temperature.
Eventually long enough to make liquids.
More electrons, more polarizable.
Bigger molecules, higher melting and boiling points.
Exist in all molecules, much weaker than other forces.
Also called Van der Waal’s forces.

12. #36

13. #38

14. #39

15. #40 (In Webassign)

16. 10.2 Liquids
Many of the properties due to internal attraction of atoms.
Beading
Surface tension
Capillary action
Stronger intermolecular forces cause each of these to increase.

17. Surface Tension Molecules at the top are only pulled inside.
Molecules in the middle are attracted in all directions.
Minimizes surface area.

18. Capillary Action Liquids spontaneously rise in a narrow tube.
Intermolecular forces are cohesive, connecting like things.
Adhesive forces connect to something else.
Glass is polar.
It attracts water molecules.

20. Beading If a polar substance is placed on a non-polar surface.
There are cohesive,
But no adhesive forces.
And Visa Versa

21. Viscosity How much a liquid resists flowing.
Large forces, more viscous.
Large molecules can get tangled up.

22. Model for Liquids
Can’t see molecules so picture them in motion but attracted to each other.
With regions arranged like solids but
with higher disorder.
with fewer holes than a gas.
Highly dynamic

23. Phases The phase of a substance is determined by three things.
The temperature.
The pressure.
The strength of intermolecular forces.
Changes of State Video

24. 10.3 Solids
Two major types.
Crystalline - have a regular arrangement of components in their structure. (table salt)
Amorphous - those with much disorder in their structure. Said to be “frozen in place”. (glass, plastic)

25. 10.4 Bonding Models for Metals
Why do metal atoms stay together? How does their bonding effect their properties?
Two Models:
Electron Sea Model: A regular array of metals in a “sea” of electrons.
Band (Molecular Orbital) Model: Electrons assumed to travel around metal crystal in MOs formed from valence atomic orbitals of metal atoms.

26. Electron Sea Model

27. 10.5 C & Si - Atomic Network Solids
Composed of strong directional covalent bonds that are best viewed as a “giant molecule”.
brittle
do not conduct heat or electricity
carbon, silicon-based
graphite, diamond, ceramics, glass
Diamond - hardest natural substance on earth, insulates both heat and electricity.
Graphite - slippery, conducts electricity.

28. Diamond - each Carbon is sp3 hybridized, connected to four other carbons.
Carbon atoms are locked into tetrahedral shape.
Strong  bonds give the huge molecule its hardness.

29. Graphite is different Each carbon is connected to three
other carbons and sp2
hybridized.
The molecule is flat with 120º angles in fused 6 member rings.
The  bonds extend above and below the plane.

30. The layers slide by each other.
This  bond overlap forms a huge  bonding network.
Electrons are free to move through out these delocalized orbitals.
The layers slide by each other.

31. 10.6 Molecular solids. S8, P4, CO2, H2O
Different molecules have different forces between them.
These forces depend on the size of the molecule.
They also depend on the strength and nature of dipole moments.
Non-Polar: Large molecules (such as I2 ) can be solids even without dipoles.
Polar: Dipole-dipole forces are generally stronger than L.D.F.
Hydrogen bonding is stronger than Dipole-dipole forces.
No matter how strong the intermolecular force, it is always much, much weaker than the forces in bonds.
Stronger forces lead to higher melting and freezing points.

32. 10.7 Ionic Solids
The extremes in dipole dipole forces-atoms are actually held together by opposite charges.
Huge melting and boiling points.
Atoms are locked in lattice so they are hard and brittle.
Every electron is accounted for so they are poor conductors - good insulators.

33. What type of solid will each substance form? #68
Diamond Quartz
PH3 NH4NO3
H2 Ar
Mg Cu
KCl C6H12O6
SF2

34. 10.8 Vapor Pressure
Vaporization - change from liquid to gas at boiling point.
Evaporation - change from liquid to gas below boiling point.
Heat (or Enthalpy) of Vaporization (Hvap) - the energy required to vaporize 1 mol at 1 atm.

35. Vaporization is an endothermic process - it requires heat.
Energy is required to overcome intermolecular forces.
Responsible for cool earth.
Why we sweat.

36. Condensation Change from gas to liquid.
Achieves a dynamic equilibrium with vaporization in a closed system.
What is a closed system?
A closed system means matter can’t go in or out.
What the heck is a “dynamic equilibrium?”

37. Dynamic equilibrium
When first sealed the molecules gradually escape the surface of the liquid.

38. Dynamic equilibrium Rate of Vaporization = Rate of Condensation
Molecules are constantly changing phase “Dynamic”
The total amount of liquid and vapor remains constant “Equilibrium”

39. Vapor pressure The pressure above the liquid at equilibrium.
Liquids with high vapor pressures evaporate easily. They are called volatile.
Increases with increasing temperature.
Decreases with increasing intermolecular forces.
Bigger molecules (bigger LDF)
More polar molecules (dipole-dipole)

40. Changes of state
The graph of temperature versus heat applied is called a heating curve.
The temperature a solid turns to a liquid is the melting point.
The energy required to accomplish this change is called the Heat (or Enthalpy) of Fusion Hfus

41. Water and Steam Water and Ice
Heating Curve for Water
Steam
Water and Steam
Water
Water and Ice
Ice

42. Heat of Vaporization Heat of Fusion
Heating Curve for Water
Slope is Heat Capacity
Heat of
Vaporization
Heat of
Fusion

43. Melting Point
Melting point is determined by the vapor pressure of the solid and the liquid.
At the melting point the vapor pressure of the solid = vapor pressure of the liquid at 1 atm

44. Boiling Point
Reached when the vapor pressure equals the pressure of the surrounding atmosphere.
Normal boiling point is the boiling point at 1 atm pressure.
Super heating - Rapid heating above the boiling point allows liquid state to exist above normal boiling point.
Supercooling - Rapid cooling below the freezing point allows liquid state to exist below the normal freezing point.

45. Practice

47. 10.9 Phase Diagrams
A plot of temperature versus pressure for a closed system, with lines to indicate where there is a phase change.
Critical temperature: temperature above which the vapor can not be liquefied.
Critical pressure: pressure required to liquefy at the critical temperature.
Critical point: critical temperature and pressure (for water, Tc = 374°C and 218 atm). Liquid & solid are indistinguishable.

48. Solid
Liquid
Gas
Critical Point
Pressure
Triple Point
Temperature

49. Solid
Liquid
Gas D D
Pressure D C
1 Atm B D A
Temperature

50. This is the phase diagram for water.
Solid
Liquid
Gas
This is the phase diagram for water.
The density of liquid water is higher than solid water.
Pressure
Temperature

51. Pressure 1 Atm Temperature This is the phase diagram for CO2
The solid is more dense than the liquid
The solid sublimes at 1 atm.
Pressure
Liquid
Solid
1 Atm
Gas
Temperature

53. 
Like most substances, bromine exists in one of the three typical phases. Br2 has a normal melting point of -7.2°C and a normal boiling point of 59°C. The triple point for Br2 is -7.3°C and 40 torr, and the critical point is 320°C and 100 atm. Using this information, sketch a phase diagram for bromine indicating the points described above.
Based on your phase diagram, order the three phases from least dense to most dense.
What is the stable phase of Br2 at room temperature and 1 atm?
Under what temperature conditions can liquid bromine never exist?
What phase changes occur as the temperature of a sample of bromine at 0.10 atm is increased from -50°C to 200°C?